1. List the first five things that come to mind when you hear this word.
Chemistry
List the first five things that come to mind when you hear this word.

2. Vocabulary: Define as many of the following words as you can:
Atomic Mass
Atomic Number
Electron
Ion
Neutron
Proton
Isotope
Valence (Electrons)
Covalent Bond
Ionic Bond
Molecule
Isomer

3. The Periodic Table

4. Three Subatomic Particles
Proton: (+) charged particle found inside the nucleus
Neutron: neutral particle found inside the nucleus
Electron: (-) charged particle found outside the nucleus in various energy levels

6. Atomic Number & Atomic Mass
Atomic Number: the number of protons (and electrons in neutral atoms)
Mass Number: the number
of protons plus the number
of neutrons (mass of nucleus)
Atomic Mass (Weight): weighted average of the isotopes for an element
Mass #
Atomic #

7. Isotopes of Hydrogen What are isotopes? Atoms of the same element with
the same # of protons, different # of neutrons

8. The Periodic Table Groups/Families (columns) Periods (rows)
Atomic Number
Atomic Mass
Valence Electrons

9. Valence Electrons
Valence Electrons: electrons in outermost energy level
Most important for bonding

10. Covalent Bonds
Valence electrons are shared between atoms

11. Sharing one or more electron pairs
Covalent Bonding
Sharing one or more electron pairs
between 2 atoms

12. Covalent Bonds

13. Ions
Atoms of the same element with a different number of electrons (charged particles)
Cations (positive ions) – lose electrons
Anions (negative ions) – gain electrons

14. CATION
“cat”ion
ca+ion
ANION
“ant”ion
n - negative

15. Ionic Bonds
Valence electrons are transferred between atoms

16. Ionic bonding

17. Ionic Bonding

18. Molecule
Smallest unit of most compounds that displays all the properties of that compound

19. Counting Atoms Na2SO4 2 - Na, 1 - S, 4 - O 7 atoms
Ca(OH) Ca, 2 - O, 2 - H
5 atoms
3 Fe2(SO3)3 6 - Fe, 9 - S, 27 - O
42 atoms

20. Isomers
Compounds with the same molecular formula but different structures (arrangement of atoms) and, therefore, different properties

21. Properties of Water
Section 2.2

22. What do you know about water?

23. Unique Properties of Water:
Most abundant compound in nearly all living things

24. Unique Properties of Water:
Density of solid water vs. liquid water
Expansion on freezing – water forms a crystalline structure that expands and is less dense than its liquid state

25. Unique Properties of Water:
Covers more than 75% of the Earth’s surface
Water 100°C or 212°F
Water 0°C or 32°F
Universal solvent

26. Unique Properties of Water:
Water has a high heat capacity
Heat capacity = the amount of heat energy required to increase its temperature, is relatively high.
Large bodies of water (oceans and lakes) can absorb large amounts of heat with only small changes in temperature.
This protects organisms living within from drastic changes in temperature.

27. What is a polar molecule?
An uneven distribution of electrons
The molecule becomes charged on each end
It means that the electrons are
NOT shared equally

28. Hydrogen Bonding
Because of their partial positive and negative charges, polar molecules such as water can attract each other.
The attraction between a hydrogen atom on one water molecule and the oxygen atom on another is known as a hydrogen bond.

29. Interaction Between Water Molecules
Negative Oxygen atomof one water molecule is attracted to the Positive Hydrogen atom of another water molecule to form a HYDROGEN BOND

31. Cohesion
Cohesion is when one water molecule sticks to another water molecule  due to hydrogen bonding
It creates surface tension (a measure of strength) at the surface of the liquid
Produces a surface “film” that allows small insects to “walk on water”

32. Surface Tension
A measure of how difficult it is to break the surface of a liquid because of hydrogen bonds.
The surface tension of a liquid results from an imbalance of intermolecular attractive forces, (cohesive forces) between molecules
A molecule in the bulk liquid experiences cohesive forces with other molecules in ALL directions.
A molecule at the surface of a liquid experiences only NET INWARD cohesive forces.

33. Adhesion Adhesion is when a water molecule sticks to another substance
(NOT another water molecule)
A piece of tape adheres to the wall
Capillary action is an example of adhesion (water moving from the roots, through the stem, & into the leaves of a flower)

34. Capillary Action
Water climbing up a small tube because of the adhesion to the sides and the cohesion between water molecules.

35. Adhesion Also Causes Water to …
Attach to a silken spider web
hold onto plant leaves

36. Cohesion and Adhesion
Mercury does not wet glass - the cohesive forces within the drops are stronger than the adhesive forces between the drops and glass. When liquid mercury is confined in a tube, its surface (meniscus) has a convex shape because the cohesive forces in liquid mercury tend to draw it into a drop.
When liquid water is confined in a tube, its surface (meniscus) has a concave shape because water adheres to the surface and creeps up the side.

37. High Heat Capacity
Water’s heat capacity, the amount of heat energy required to increase its temperature, is relatively high.
Because of the multiple hydrogen bonds between water molecules, it takes a large amount of heat energy to cause those molecules to move faster and raise the temperature of the water.
Large bodies of water, such as oceans and lakes, can absorb large amounts of heat with only small changes in temperature. This protects organisms living within from drastic changes in temperature.

38. Universal Solvent Hydrophilic Hydrophobic
A substance that is attracted to water (a polar molecule)
Examples: sugar, alcohol
Hydrophobic
A substance that repels (is afraid of) water (a nonpolar molecule)
Examples: oil
The Chemistry of Water

39. Solutions and Suspensions
Mixtures
Solutions and Suspensions

40. What is a Mixture?
Substance composed of two or more elements or compounds that are mixed together, but NOT chemically combined.
Mixture Examples: salt & pepper; sugar & sand; salt & sand; Earth’s atmosphere; soil

41. Solutions and Suspensions
Mixture – a material that consists of 2 or more elements or compounds physically mixed together but are not chemically combined.

42. Homogeneous Mixture (Solution)
Homogeneous mixture in which one substance is dissolved in another
Examples: sugar in water; salt in water; hot chocolate; tea, Kool-aid
Salt Dissolving in Water

43. Parts of a Solution All solutions contain a solute and a solvent.
Solvent is the substance doing the dissolving.
Example: water
Solute is the substance that is being dissolved.
Example: salt

44. Heterogeneous Mixture (Suspensions)
Mixture of water and non-dissolved materials (heterogeneous)
Examples: blood; oil and water

45. Application: 1 – Choose one of your examples of a
solution. Explain how the parts
(components of the solution) can be
separated.
2 – For the same example, identify the
solute and solvent in the solution.
Explain your reasoning.

46. Example of a Solution: Water is the universal solvent.
The solute (sugar) dissolves in the solvent (water).
The solute (Kool-aid) dissolves in the solvent (water).
The solute (oxygen) dissolves in the solvent (nitrogen).
Water is the universal solvent.

47. Homeostasis a. Makes a good insulator b. Resists temperature change
Ability to maintain a steady state despite changing conditions
Water is important to this process because…
a. Makes a good insulator
b. Resists temperature change
c. Universal solvent
d. Coolant
e. Ice protects against temperature extremes (insulates frozen lakes) 4

48. Acids, Bases, pH, & Buffers
All about the pH Scale

49. The Water Molecule:
Water dissociation animation

50. Acids Compounds that release H+ into solution H+ = hydrogen ion
If [H+] > [OH-] then the
solution is acidic.
HCl  H+ + Cl-

51. Bases If [OH-] > [H+] then the solution is basic
Compounds that release
OH- into solution
OH- = hydroxide ion
If [OH-] > [H+] then the
solution is basic
NaOH  Na+ + OH-

52. The pH scale (power of Hydrogen) as in Hydrogen ions
Measures the relative concentration of H+ and OH-

53. Interpreting the pH Scale
pH is based on a logarithmic scale
A pH of 5 is 10× more concentrated than a pH of 6 (in terms of H+ concentration)
A pH of 9 is 100× more concentrated than a pH of 11 (in terms of H+ concentration)

54. Classifying Acids & Bases
Strong acids have a pH value between 1-3
Weak acids have a pH value between 4-6
Bases:
Strong bases have a pH value between 12-14
Weak bases have a pH value between 8-11

55. Buffers
Weak acids or bases that can react with strong acids or bases to prevent sharp sudden changes in pH
Produced naturally by the body to maintain homeostasis
The pH value in most cells is
The pH of stomach acid is 2
The pH of the blood is 7.4

56. Blood Buffer – pH needs to be 7.4
Double arrows mean
reactions are in
equilibrium
bicarbonate ion (buffer)
(Kidney – carried
away in urine)
Free hydrogen ions
Carbonic Acid
(Lungs – carried
away as CO2)

57. Blood Buffer – pH needs to be 7.4
You can see that if the reactions go to the right, hydrogen ions are released, making the solution more acidic, and if the reactions go to the left, hydrogen ions are sucked up, making the solution more basic.  The question that is probably eating away at you right now is:  "how do I know which way the reactions are going to go-- to the right or to the left?“
Good question!  The bicarbonate ion isn't a very strong acid or base.  It doesn't have to go one way or the other all the time.   Instead, the direction it goes depends on the solution it is in.  You see, if it is in an acidic solution, there are lots of hydrogen ions floating around; in this situation, the presence of tons of hydrogen ions will force the reaction to go to the left.  Do you see why?  Look at the reaction again.  All you need for the reaction to be able to go to the left is available hydrogen ions (see the H+ in the middle that is needed for the reaction to go to the left?).  So if bicarbonate ions find themselves in acidic solutions, they tend to act like the weak base they can be, and suck up the excess hydrogen ions.
Consider what will happen if bicarbonate is put into a basic solution. A basic solution has very few hydrogen ions floating around.  In this condition, the bicarbonate reaction cannot proceed to the left, since no hydrogen ions are available.  Instead, it will go to the right, producing hydrogen ions.  In this manner, bicarbonate ions will tend to act like the weak acid they can be, and release hydrogen ions.

58. Use the structure of a water molecule to explain why it is polar.
Why is water such a good solvent?
What is the difference between a solution and a suspension?
How do hydrogen bonds between water molecules occur?
How does the body counteract a drop in blood pH during strenuous exercise?
If a strong acid is dissolved in pure water will the pH of the solution be greater or less than 7?

59. Water Chemistry
Please draw the structure of a water molecule (H2O).
Label the charges on the molecule that make it polar (negative and positive ends).
See if you can connect the water molecule that you drew to two other molecules of water using hydrogen bonds.
Can you think of any unique properties that water displays?

60. Water Chemistry
Explain why water is “polar” in your own words.
Explain capillary action in your own words.
Explain surface tension in your own words.

61. According to the scale, what substance is neutral?
What solution has an [H+] of 0.01?
Which substance produces the most OH- ions in solution?
How many more times basic is pH 14 compared to pH 11.

62. Chemical Reactions and Enzymes
Section 2.4

63. Objectives Distinguish between potential and kinetic energy.
Differentiate between reactants and products in a chemical reaction.
Explain the role of a catalyst in chemical reactions.

64. What is energy? The ability to do work or cause a change.
Can exist in different forms:
thermal radiant
chemical nuclear
mechanical electrical

65. Forms of Energy: POTENTIAL ENERGY  stored energy
KINETIC ENERGY  energy of motion
FREE ENERGY  the energy available to do work

66. Activation Energy
the energy needed to start a chemical reaction.

67. Chemical Reactions:
Everything that happens in an organism is based on chemical reactions.
A chemical reaction is a process that changes one set of chemicals into another set of chemicals.
CO2 + H2O  C6H12O6 + O2

68. Chemical Reactions:
The elements or compounds that enter into the reaction are the reactants.
The elements or compounds produced by the reaction are the products
Bonds are broken/formed in chemical reactions.

69. What are the Reactants? Products?
CO2 + H2O  C6H12O6 + O2
Reactants
Products
 means “Yields” or “Produces”

70. Energy RELEASED Reactions
Activation Energy: the energy needed to start a chemical reaction.
Exergonic (Exothermic) Reactions - involve a net release of free energy
Often happen spontaneously

71. Energy ABSORBED Reactions
Activation Energy: the energy needed to start a chemical reaction.
Endergonic (Endothermic) Reactions - involve a net absorption of free energy
Will not occur without a source of energy

72. Draw the enzyme pathway on your graph
Enzymes
Enzymes – are proteins that act as a biological catalyst.
Catalyst – a substance that speeds up a chemical reaction by lowering the activation energy.
Draw the enzyme pathway on your graph

73. Enzyme Action
Substrate – the reactants of an enzyme catalyzed reaction
Active Site – where the substrate and enzyme join.
Have complimentary shapes
Can only bind with a specific molecule – “induced fit”

74. How Enzymes Work: Enzymes act on a specific substrate.
(Shapes fit together like a lock and a key)
A small area on the enzyme, called the active site, can attract and hold only a specific substrate.
The enzyme acts as a biological catalyst, which accelerates the rate of the chemical reaction.
The enzyme reduces the activation energy needed by weakening the chemical bonds in the substrate.
The enzyme is then released unchanged.

75. Effects of Enzymes:

76. Regulation of Enzyme Activity
Enzymes work best at certain pH levels and temperatures.
Most enzymes in humans work best at 37C
Denaturation – a change in the enzymes shape
Enzyme becomes nonfunctional
Heat and pH can cause denaturation
Some enzymes can be turned on and off
Lactose intolerance

78. What letter represents the product?
What letter represents the reactant?
What letter represents the activation energy?
How should the x and y axis be labeled?
Is the reaction exergonic or endergonic? Explain